EXCHANGE 


The  Formation  of  Addition  Compounds  Between  Formic 

Acid  and  Metallic  Formates.    A  Discussion  of 

the  Factors  Affecting  the  Stability  of 

These  Compounds. 


by 

HOWARD  ADLER,  B.S.,  M.A. 


DISSERTATION  ONIV 

Submitted  in  Partial  Fulfillment  of  the  Requirements 

for  the  Degree  of  Doctor  of  Philosophy  in 

the  Faculty  of  Pure  Science  of 

Columbia  University. 


NEW  YORK  CITY 
1920 


The  Formation  of  Addition  Compounds  Between  Formic 

Acid  and  Metallic  Formates.    A  Discussion  of 

the  Factors  Affecting  the  Stability  of 

These  Compounds. 


by 

HOWARD  ADLER,  B,S.,  M.A. 

tt 


DISSERTATION 

Submitted  in  Partial  Fulfillment  of  the  Requirements 

for  the  Degree  of  Doctor  of  Philosophy  in 

the  Faculty  of  Pure  Science  of 

Columbia  University. 


NEW  YORK  CITY 
1920 


ACKNOWLEDGMENT 

The  author  wishes  to  express  to  Professor  James 
Kendall  his  sincere  gratitude  for  the  suggestion  of 
the  problem,  and  for  his  helpful  advice  throughout  the 
course  of  the  investigation. 

The  author  also  wishes  to  acknowledge  with 
thanks  the  co-operation  of  the  other  members  of  the 
Chemistry  Department  of  Columbia  University. 


451724 


ABSTRACT  OF  DISSERTATION 

1.  What  was  attempted? 

2.  In  how  far  were  the  attempts  successful? 

3.  What  contribution  actually  new  to  the  science  of  chemistry 

has  been  made  ? 

1.  The  attempt  was  made  to  demonstrate  the  applicability 
of  rules  previously  formulated  governing  the  variation  of  ad- 
dition compound  formation,  to  systems  of  the  type  HX-RX. 
The  particular  systems  studied  were  the  formic  acid— metallic 
formate  series.    A  few  cases  of  the  series  acetic  acid — metallic 
acetate  were  also  studied. 

2.  (a)   It  has  been  shown  that  the  extent  of  combination 
between  formic  acid  and  the  metallic  formates  varies  uniformly 
with  the  position  of  the  metal  (R)  in  the  electromotive  series. 
As  the  position  of  R  changes   from  potassium  through  the 
series,  the  extent  of  compound  formation  decreases  to  a  mini- 
mum (in  the  neighborhood  of  hydrogen).    A  similar  variation 
occurs  in  the  acetate  systems,  with  a  slight  increase  in  the  ex- 
tent of  combination  when  the  position  of  R  is  considerably  be- 
low hydrogen. 

(b)  It  has  also  been  shown  that  the  variation  in  the  extent 
cf  compound  formation  is  parallel  to  the  change  in  conductivity. 

(c)  The  change  in  solubility  follows  regularly  the  change 
in  compound  formation. 

3.  (a)    The  initial  steps  in  the  formulation  of  a  general- 
ized theory  of  ionization  have  received  additional  experimental 
confirmation. 

(b)  Further  development  of  methods  for  determining  the 
existence  and  extent  of  compound  formation  in  solution  has 
been  indicated. 

(c)  Evidence  has  been  obtained  of  causal  relationships 
between  compound  formation,  conductivity,  solubility,  and  di- 
versity of  components  in  systems  of  the  general  type  HX — RX. 

(d)  In   the  course  of  the  investigation,   five  new  acid 
formates  have  been  isolated. 


THE  FORMATION  OF  ADDITION  COMPOUNDS  BE- 
TWEEN  FORMIC  ACID  AND  METALLIC   FOR- 
MATES.   A  DISCUSSION  OF  THE  FACTORS 
AFFECTING  THE  STABILITY  OF 
THESE  COMPOUNDS. 


INTRODUCTION 

In  the  preceding  articles  of  this  series1,  it  has  been  con- 
clusively demonstrated  that  the  extent  of  compound  formation 
in  a  solution  depends  essentially  upon  the  difference  in  charac- 
ter of  the  two  components.  On  the  basis  of  the  generalization2 
that  compounds  increase  in  stability  uniformly  with  increas- 
ing divergence  in  the  basic  or  acidic  nature  of  the  two  compo- 
nents, it  has  been  possible  to  predict  the  relative  extents  of 
combination  in  various  systems,  as  well  as  the  relative  stability 
of  the  complexes  formed.  Agreement  of  experimental  results 
with  these  predictions  has  been  extremely  satisfactory3. 

The  study  of  aqueous  systems  of  the  type  HX-HOH,  and 
to  a  less  extent,  of  the  type  ROH-HOH4,  revealed  the  fact  that 
the  generalization  developed  for  non-aqueous  solvents  could  be 
applied,  with  equally  good  results,  to  solutions  in  water.  Prac- 
tically complete  predetermination  of  the  extent  of  the  reaction 
HX  +  HOH  ^  HX-HOH  from  left  to  right  was  possible. 

The  results  of  the  investigation  mentioned  above  led  to  the 
correlation  of  ionization  and  compound  formation5.  It  was 
shown  that  for  systems  of  the  type  HX-HOH,  and  ROH-HOH, 
the  extent  of  ionization  varied  directly  with  the  extent  of  com- 
bination between  the  two  components  of  the  system.  The 
causal  relationship  between  the  two  phenomena,  which  this 
variation  suggested,  has  been  verified  by  Gross0.  The  conduc- 


1  Kendall  and  Booge,  J.  A.  C.  S.  38,  1712  (1916).    For  resume  see  Ken- 

dall, Booge  and  Andrews,  ibid,  39,  2304  (1917). 

2  Lewis,  System  of  Physical  Chemistry,  V.  I,  pg.  417. 

3  The  exceptional  behavior  of  phenols  has  been  noted.     See  Kendall, 

Booge  and  Andrews-  loc.  cit.  p.  2306;  also  Kendall  J.  A.  C.  S.  38. 
1317,  1322  (1916) ;  36,  1240  (1914). 

4  Kendall,  Booge  and  Andrews,  loc.  cit. 

5  Kendall  and  Booge,  J.  A.  C.  S.  39,  2323  (1917).    - 

6  Gross,  Columbia  University  Dissertation,  1919. 


tivity  measurements  made  by  this  investigator  confirm  the 
hypothesis  that  ionization  in  solution  is  preceded  by  compound 
formation  between  solvent  and  solute. 

The  generalization  concerning  compound  formation  was 
postulated  for  aqueous  systems  of  the  type  ROH-HOH  in  the 
following  form1 : 

(a)  No  indication  of  hydrate  formation  with  extremely 
weak  bases. 

(b)  A  regular  increase  in  the  extent  of  combination  in 
the  liquid  state  as  the  strength  of  the  base  is  increased. 

(c)  Extensive  compound  formation  with  transition  and 
strong  bases. 

(d)  Increase  in  complexity2  as  well  as  stability  of  hydrates 
with  the  strength  of  the  base. 

Experimental  verification  of  these  statements  for  solutions 
of  the  type  water-base  (as  afforded  by  the  work  of  previous 
investigators)  was  presented  at  the  time  of  their  proposal.  It 
is  to  an  extension  of  this  topic  that  the  present  paper  is  de- 
voted. Additional  data  upon  this  problem  are  essential  to  a 
more  general  formulation  of  the  ionic  theory,  since,  as  has 
already  been  mentioned3,  more  complete  knowledge  of  the  com- 
paratively simple  systems  RX-HX  and  ROH-RX  is  necessary 
before  any  quantitative  knowledge  of  the  complex  system 
RX-H2O  can  be  gained. 

Solvent-Base-Systems.  With  a  generalization  of  the 
theory  of  solutions  will  come,  of  necessity,  a  broadening  of 
definitions,  so  as  to  remove  the  restrictions  now  imposed  by 
the  limited  range  of  applicability  of  the  accepted  theory.  In 
terms  of  the  broader  conceptions,  a  base  is  defined  as  a  binary 
compound,  which  on  solution  yields  the  same  negative  ion  as 
the  negative  radical  of  the  solvent4.  With  this  in  mind,  it  is 


1  Kendall,  Booge  and  Andrews,  loc.  cit.  p.  2320. 

2i.  e.    compounds  of  the  type   (ROH)a(HOH)b  can  be  isolated  when 
there  is  considerable  divergence  between  the  components. 

3  Kendall,  Booge  and  Andrews,  loc.  cit,  p.  2322. 

4  This  terminology  has  been  used  by  Schlesinger  and  Calvert,  T.  A.  C.  S. 

33,  1933  (1911"). 


obvious  that  the  non-aqueous  system  HX-RX  is  identical  in 
nature  with  the  system  ROH-HOH.  Hence  it  is  of  interest  to 
see  whether  the  postulates  proposed  for  the  latter  type  will 
holdi  equally  well  for  the  former.  This  condition  is  essential 
to  the  evolution  of  a  more  comprehensive  theory  of  solutions. 

Criterion  of  Diversity.  Since  diversity  of  the  two  compo- 
nents is  postulated  as  necessary  for  combination  between  them, 
ir  becomes  necessary  to  arrive  at  some  criterion  of  the  extent 
of  divergence  in  systems  of  the  type  HX-RX.  The  X  radical 
being  common  to  both  components,  the  "difference"  between 
H  and  R  must  afford  a  measure  of  the  tendency  to  complex 
formation,  i.  e.  of  the  tendency  of  the  reaction  HX  +  RX  ^ 
HX-RX  to  go  forward.  The  best  available  criterion  of  this 
divergence  is  the  relative  positions  of  the  metal  (R)  and  hydro- 
gen in  the  electromotive  series.  This  follows  from  the  rela- 
tionship which  exists  between  the  electromotive  series  (or 
electrode  potential  series,  since  these  are  identical  in  order)  and 
the  chemical  activity  of  the  metals1.  The  relative  activities  of 
the  metals  are  given  by  the  order  of  the  electromotive  series, 
the. activity  decreasing  in  order  from  K  down  the  series  to 
the  noble  metals. 

It  is  to  be  expected,  then,  that  the  higher  the  position  of 
the  metal  (R)  in  the  series,  the  stronger  will  be  the  resulting 
base  (RX)2.  As  R  is  varied,  and  approaches  hydrogen,  the 
strength  of  the  resulting  base,  as  evidenced  by  the  extent  of 
compound  formation  and,  consequently3,  of  ionization,  should 
diminish.  In  the  immediate  neighborhood  of  hydrogen,  the 
base  should  be  very  weak.  On  continuing  the  variation  be- 
yond hydrogen,  the  difference  between  the  components  be- 
comes more  pronounced,  the  lower  the  position  of  the  metal  in 
the  series.  This  divergence  should  result  in  the  formation  of 


1  For  a  full   discussion  of  chemical   affinity  and   its   measurement  by 

e.  m.  f.,  see  particularly  Lewis,  System  Physical  Chemistry,  V.  2, 
Ch.  XII.  Also  Lehfeldt,  Electrochem.,  pp.  181,  182,  194.  Mellor, 
Modern  Inorganic  Chem.,  361-376;  LeBlanc,  Electrochem.,  267  et  seq. 

2  This  ought  to  be  equally  true,  irrespective  of  the  nature  of  X — i.e., 

whether  it  be  NO3,  SO4,  OH— etc. 

3  Kendall  and  Booge,  loc.  cit.,  p.  2324. 

8 


bases  stronger  than  those  of  the  metals  located  near  hydrogen 
in  the  series1. 

It  is  seen  that  on  the  basis  of  the  above  assumption  the 
strength  of  the  base  RX  should  diminish  to  a  minimum  and 
then  increase  again,  as  R  is  varied  from  one  extreme  of  the 
electromotive  series  to  the  other.  Those  properties  which  are 
usually  associated  with  the  term  "strength" — extent  of  ioniza- 
tion  (compound  formation),  and,  as  will  be  seen  from  the 
sequel,  solubility,  should  undergo  concomitant  variation.  The 
experimental  work  to  be  described,  will  seek  to  establish  the 
validity  of  this  argument. 

SYSTEMS  HX  RX.    ACID  SALTS. 

There  are  in  the  literature  numerous  references  to  com- 
pounds of  the  type  (RX)a  (HX)b,  i.  e.,  acid  salts.  There  have 
been,  however,  very  few  systematic  investigations  of  such  salts, 
with  a  view  to  the  correlation  of  the  fact  of  their  existence 
with  theory. 

It  might  be  worth  considering  briefly  those  compounds 
mentioned  in  the  literature,  in  order  to  see  how  far  they  are  in 
agreement  with  the  requirement  of  the  theoretical  basis  de- 
veloped above.  A  complete  survey  would  not  be  of  any  value, 
because  such  acid  salts  as  bisulfites,  bicarbonates,  and  others, 
could  never  lend  themselves  to  systematic  study,  since  the  acids 
exist  only  in  solution.  The  following  review  will  be  limited 
to  those  cases  where  the  compounds  can  be  obtained  from  the 
pure  acid  and  base. 


1  The  position  of  the  metals  in  the  electromotive  series  (or  rather,  elec- 
tro-affinity) as  a  factor  in  determining  the  properties,  of  metallic 
compounds,  such  as  the  chlorides,  has  been  used  by  Bodlander,  Z. 
Phys.  Chem.  27,  55  (1898),  and  Abegg  and  Bodlander,  Z.  Anorg. 
Chem.  20,  453  (1899).  They  attempted  to  show  for  example,  that 
the  solubility  of  the  chlorides  increased  as  the  position  of  the  metal 
varied  from  bottom  to  top  of  the  series,  i.  e.  from  Ag  to  K.  This  is 
notoriously  not  the  case ;  which  fact  led  to  the  use  of  additional 
hypotheses  to  maintain  the  original  thesis.  The  procedure  was  not 
altogether  warranted.  The  failure  to  establish  the  validity  of  the 
propositions  advanced  is  probably  due  to  the  total  neglect  of  the 
influence  of  the  solvent.  Furthermore,  the  systems  examined  were 
of  the  complex  type  RY-HX  or  RY-H0O,  and  the  complications 
introduced  by  the  fourth  radical  prevented  any  real  connection  be- 
tween the  electromotive  series  and  properties  such  as  solubility  from 
being  discovered. 


The  only  acid  whose  acid  salts  have  been  completely  ex- 
amined is  sulfuric  acid1.  The  results  of  the  entire  investigation 
are  to  be  published  shortly,  so  that  a  complete  discussion  of 
this  system  is  not  now  advisable.  The  results  in  general  are, 
however,  in  highly  satisfactory  agreement  with  the  theory. 
There  has  been  no  other  complete  series  studied,  with  any 
theoretical  objective.  The  compounds  mentioned  in  the  liter- 
ature, as  conditioned  above,  are  given  in  Table  I,  which  follows  : 

TABLE  I. 

Solvent        Base  Compounds 

HN03     KN03  KN03-2HN03;  KNO3-HNO32 

NH4NO3      NH4NO3-2HNO3;  NH4NO3-HNO32 
HAc        KAc  KAc  2HAc3;  KAc-HAc4 

NaAc  NaAc-2HAc;  NaAc-H.V 

NH4Ac         NH4Ac-HAc6 
Li  Ac  LiAc-HAc7 

TIAc  TIAc-HAc7 

HF         KF  KF-3HF8;  KF-HF9 

NaF  NaF-HF10 

LiF  LiF-HF11 

NH4F          NH4F-HF12 
AgF  AgF-3HF13;  AgF-HF13 

The  existence  of  all  of  these  compounds,  as  well  as  that 
of  the  hydrates  of  bases  which  have  already  been  enumerated14, 
is  in  agreement  with  the  requirements  of  the  theory.  All 


JLandon,  J.  A.  C.  S.,  42,  2131  (1920).    Completed  by  Davidson,  as  yet 

unpublished. 
2Groschuff,  Ber.  _37,  1488  (1904).    The  same  author's  work  on  formic 

acid  will  be  referred  to  in  the  following  section. 
sLescoeur,  Ann.  Chim.  Phys.  (6),  28,  245  (1893). 
*Melsens,  Compt.  Rend.  19,  611  (1844). 
5  Lescoeur,  loc.  cit.,  pg.  241. 
6Reik,  Monatshefte  f.  Chemie,  23,  1033  (1902). 
7  Lescoeur,  Bull.  Soc.  Chim.,  24,  517  (1875). 
s  Moissan,  Compt.  Rend.,  106,  547  (1888). 
9Abegg,  Hand.  Anorg.  Chem.,  2-1,  343. 
10Abegg,  ibid.,  220-221. 

11  Ibid.,  pg.  120. 

12  Marignac,  Ann.  Min.  (5)  15,_221  (1859). 

«  Guntz,  Bull.  Soc.  Chim.  (3)  U  114  (1895). 

14  Kendall,'  Booge  and  Andrews,  loc.  cit.,  p.  2320. 

10 


metals  whose  "bases"  form  addition  compounds  of  the  acid- 
salt  type,  are  either  strongly  electropositive  or  strongly  elec- 
tronegative. It  is  especially  noteworthy  that  silver  fluoride  is 
soluble  in  hydrofluoric  acid,  and  is  extensively  solvated. 

It  is  evident  that  the  data  existing  are  not  sufficient  to 
supply  a  rigorous  test  of  the  validity  of  the  argument.  The 
varying  reliability  of  results  from  scattered  sources,  the  lack 
of  completeness  in  all  the  series  examined,  as  well  as  the  ab- 
sence of  investigations  of  freezing-points  so  as  to  permit  the 
determination  of  relative  extents  of  combination  throughout 
the  series,  all  tend  to  diminish  the  value  of  any  conclusion  which 
may  be  drawn.  For  these  reasons,  it  was  deemed  necessary  to 
determine  as  completely  as  possible  the  freezing-point  curves 
for  the  series  formate-formic  acid.  In  this  series  the  only  part 
which  was  varied  was  that  which  corresponds  to  R  in  the  type 
system  RX — HX.  After  the  effect  of  this  variation  of  R  has 
been  determined,  the  role  of  X  can  be  more  exactly  examined, 
by  a  study  of  several  series  similar  to  the  one  in  question. 

FORMIC  ACID  AS  SOLVENT 

It  has  already  been  determined  by  Schlesinger  and  col- 
laborators1, that  solutions  of  the  formates  in  formic  acid  are 
excellent  conductors.  The  alkali  formates  are  highly  ionized, 
and  are  entirely  analogous  to  the  alkali  hydroxides  in  water. 
It  has  also  been  shown2  that  the  conductivity  of  a  solution  de- 
pends upon  two  factors — (a)  the  extent  of  compound  forma- 
tion and  (b)  the  extent  of  dissociation  of  the  complexes  into 
ions  of  opposite  charge.  Since  the  metallic  formates  form  high 
ly  conducting  solutions  in  formic  acid,  it  follows  that  they  are 
highly  solvated  in  solution. 

The  formates  in  formic  acid  should  give  rise  to  compound 
formation,  varying  in  extent  with  the  position  of  the  metal  in 
the  electromotive  series3.  The  variation  in  compound  forma- 


1  With  Calvert  J.  A.  C.  S.  33,  1924  (1911)  ;  Martin,  ibid,  36,  15S9  (1914) ; 

Coleman,  ibid,  38,  271  (1916)  ;  Mullinix,  ibid,  41,  72  (1919)  ;  Reed,  41_, 
1921  (1919). 

2  Kendall  and  Booge,  loc.  cit,  p.  2324  (1917).    Gross,  loc.  cit,  pg.  7. 

3  The  order  of  increasing  conductivity  is  given  as  that  of  increasing 

electrolytic   solution   tension,    Schlesinger   and   Coleman,    loc.    cit., 
p.  278. 

11 


tion  should  parallel  the  change  in  conductivity.  To  test  the 
validity  of  these  conclusions,  a  representative  series  of  the  for- 
mates was  examined,  namely,  K,  Na,  Li,  NH4,  Ba,  Ca, 
Mg  •  Zn,  Ni,  Pb,  Cu  and  Ag.  The  solubilities  of  these  for- 
mates were  determined,  using  the  freezing-point  method,  as 
described  below. 

Due  to  unavoidable  complications,  inherent  in  the  nature 
cf  the  solvent  (i.  e.  of  X),  it  was  ndt  possible  to  work  with 
silver.  In  order  that  the  increase  of  compound  formation  as 
R  is  varied  below  H  might  be  demonstrated,  several  members 
of  the  acetate-acetic  acid  series  were  studied.  Na,  Zn,  Ni, 
Fe  (ic)  and  Ag  were  taken  as  representing  the  different  por- 
tions of  the  electromotive  series. 

The  agreement  between  the  deductions  from  the  theoreti- 
cal considerations  discussed,  and  the  results  of  these  experi- 
ments ought  to  furnish  sufficient  evidence  to  demonstrate  the 
applicability  of  the  generalization  given  above  to  systems  of 
the  type  HX— RX. 

EXPERIMENTAL 

EXPERIMENTAL  PROCEDURE 

Freezing-point  curves  for  mixtures  of  formate  and  formic 
acid  were  determined  in  the  usual  manner1.  Points  on  the 
curves  were  taken  at  intervals  of  from  2  to  3  molecular  % ; 
at  points  of  change  of  phase,  the  intervals  were  small  enough 
to  fix  accurately  the  different  branches  of  the  curve.  Each 
point  was  determined  at  least  twice. 

Those  mixtures  which  gave  a  melting-point2  below  sixty 
degrees  (60°C)  were  investigated  in  freezing  point  tubes  by 


1  See  Kendall  and  Booge,  J.  A.  C.  S.  38,  1718  (1916),  and  Landon,  loc. 

cit,  for  discussion  of  method.  In  some  cases,  i.  e.,  at  low  tempera- 
tures, considerable  supercooling  was  encountered.  The  mixtures 
were  then  cooled  in  CO2 — acetone  paste  and  allowed  to  warm  up 
slowly,  with  stirring,  to  induce  crystallization. 

2  The  temperatures  which  follow  refer  to  the  point  at  which  a  negli- 

gibly small  amount  of  the  solid  phase  is  in  equilibrium  with  the 
solution. 

12 


the  usual  method.  Every  precaution  was  exercised  to  insure, 
as  far  as  possible,  anhydrous  conditions.  The  addition  of 
formate  to  the  acid  was  done  with  the  aid  of  a  specially  de- 
signed weighing  bottle1,  thus  reducing  exposure  to  a  minimum. 
The  stirrer  was  connected  to  the  stopper  by  means  of  rubber 
tubing,  the  system  being  in  this  way  entirely  closed.  .The  com- 
position given  for  any  of  these  solutions!  is  accurate  within 
less  than  ±0.05%.  Above  about  60°C,  recourse  was  had  to 
sealed  bulbs2,  because  of  the  increasing  vapor  pressure  of  the 
formic  acid.  These  bulbs  were  so  blown  as  to  reduce  the  air 
space  to  a  minimum,  thus  decreasing  the  amount  of  solvent 
present  as  vapor  to  a  negligible  magnitude.  The  composition 
given  for  solutions  whose  freezing-points  were  measured  in 
bulbs,  may  be  taken  as  accurate  within  —  0.1  molecular  per  cent. 

The  bath  in  which  the  tube  or  bulb  was  placed  during  the 
determination  of  the,  melting  point,  varied  with  the  tempera- 
ture range  in  which  the  point  lay.  Those  baths  used,  and  the 
temperature  interval  of  their  use,  were : 

Acetone +  CO2  (solid)  Up  to  — 25°C. 

HNO3  +  ice  —25°  — 0° 

NaCl  +  ice  -15°— 0° 

Water  0°  —100°. 

70  mol.  %  H2SO4;   30  mol.  %(NH4)2SO4  above  100°. 

Considerable  attention  was  paid  to  the  factors  affecting 
thermal  equilibrium  between  the  tube  or  bulb  and  the  bath. 
This  resulted  in  the  following  precautions  being  observed,  to 
avoid  any  appreciable  error  from  this  cause : 

1.     The  tube  and  bath  were  stirred  constantly. 


1  See  Landon,  Columbia  University  Dissertation  (1920),  pg.  9,  for  de- 

tailed description  and  cut  of  this  bottle. 

2  Points   in  bulbs  were,  of  course,  determined  under  excess   pressure, 

i.  e.,  the  vapor  pressure  of  the  system  plus  the  pressure  of  the  en- 
closed air.  Since  the  limiting  temperature  was  160°C,  and  the  effect 
of  pressure  on  the  freezing  points  so  very  small  (probably  <C0.01° 
per  atmosphere),  there  \vould  be  no  advantage  in  attempting,  if  it 
.  were  possible,  to  reduce  all  freezing-points  to  atmospheric  pressure. 

13 


2.  The  temperature  was  changed  slowly  enough  so  as 
to  maintain1,   as   nearly   as  practicable,   thermal   equilibrium 
throughout  the  heating  process. 

3.  The  bulbs  were  sufficiently  thin  to  prevent  lag  when 
the  above  precautions  were  observed2. 

The  effects  of  draughts,  and  radiation  at  higher  tempera- 
tures, were  excluded  by  the  use  of  an  asbestos  shield  sur- 
rounding1 the  bath.  This  shield  had  glass  windows  to  permit 
observation  of  the  bath. 

TEMPERATURE   MEASUREMENT 

Temperatures  were  measured  by  means  of  three  mercury 
thermometers,  graduated  in  tenths  of  a  degree  (C.),  and  having 
the  respective  ranges,  —  35°  —  +  25° ;  0°  —  100° ;  100°  —  200°. 
These  were  calibrated  at  0°  and  100°,  and  the  two  with  lower 
range  were  compared  with  a  certified  thermometer  at  inter- 
mediate points.  The  100° — 200°  thermometer  was  tested  at  the 
boiling  points  of  pure  monobrombenzene  and  aniline  (Kahl- 
baum),  giving  results  in  agreement  with  the  literature.  Hence 
it  was  considered  as  correct  within  the  limits  of  experimental 
error  (as  discussed  below). 

The  correction  for  exposed  stem  was  determined  experi- 
mentally3 by  measuring  constant  temperatures4  with  the  ther- 
mometer exposed,  and  then  repeating  with  the  thread  entirely 
immersed.  The  length  of  thread  exposed  was  the  same  as 
would  be  left  outside  the  bath  in  a  determination  of  a  melting 
point.  The  corrections  obtained  were  plotted  against  tempera- 


1  App.  0.2°  per  minute  was  the  average  rate  of  heating.    This  was  varied 

slightly,  according  as  the  slope  of  the  curve  changed.  Where  the 
rate  of  change  of  composition  with  temperature  was  high,  the  rate 
of  heating  was  diminished.  The  rate  was  increased  slightly  under 
the  reverse  circumstances. 

2  This  is  proven  by  concordant  results  when  points  on  the  same  portion 

of  a  curve  were  determined  by  either  method. 

3  The  apparatus  used  for  this  determination  consisted  of  a  glass  tube, 

resembling  the  outer  jacket  of  a  Victor  Meyer  apparatus,  to  the  top 
of  which  a  Liebig  condenser  was  attached.  Its  use  was  suggested 
by  Dr.  P.  M.  Gross. 

4  Boiling  liquids  of  good  quality,  not  necessarily  ultra  pure,  give  tem- 

peratures constant  to  0.1°  for  a  sufficient  length  of  time  to  permit 
both  measurements  being  taken.  . 

14 


tures,  and  from  the  graph  so  made,  the  stem  correction  for  any 
intermediate  point  could  be  read1. 

X:  Precision  of  Measurements — The  freezing-point  of  a  mix- 
ture prepared  by  the  above  method  and  determined  with  one 
of  the  thermometers  just  described,  possesses  a  definite  pre- 
cision value.  This  depends  not  only  upon  the  temperature 
interval  in  which  the  point  lies,  but  also  upon  the  nature  of  the 
curve.  This  is  due  to  the  fact  that  it  is  easier  to  determine, 
with  any  desired  precision,  a  point  which  lies  on  a  flat  curve 
than  one  which  is  on  a  steep  curve2. 

The  fact  that  both  of  these  factors,  temperature  and  slope, 
have  to  be  considered  makes  it  difficult  to  give  any  definite 
probable  values.  Those  which  follow  are  to  be  taken  as  ap- 
proximations, true  for  the  average  type  of  curve  only3 : 

Temperature  Interval  Possible  Error 

_35«  to  _IQO  ±  o.2°  to  ±  0.5° 

-10°  to  +  100°  ±0.1°  to  ±0.2° 

100°  to      200°  ±  0.2°  to  ±  0.5° 

ANALYSIS  OF  COMPOUNDS 

The  composition  of  the  solid  phase  separating,  in  every 
case  where  it  was  not  evident  from  the  curve,  was  determined 
by  analysis.  A  mixture  of  suitable  composition  was  prepared, 
and  the  substances  to  be  analyzed  frozen  out.  The  compound 
was  then  collected  in  a  Gooch  crucible,  the  solution  being  drawn 
through  by  suction.  The  filtration  was  carried  out  under  anhy- 
drous conditions4.  The  solvent  adhering  to  the  crystals  was  re- 
moved by  sucking  air  dried  by  CaCl2  through  for  sufficient 
time  to  guarantee  complete  removal. 


1  While  it  is  evident  that  these  corrections  hold  exactly  only  when  the 

temperature  surrounding  the  stem  is  the  same  as  when  the  correc- 
tions were  determined,  a  change  of*3°  or  4°  in  room  temperature 
produces  no  appreciable  effect  upon  the  values. 

2  Flat  and  steep  refer  to  the  slopes  of  the  curves — the  change  of  tem- 

perature with  respect  to  a  slight  change  in  composition  is  small 
and  large  on  the  respective  curves. 

3  As  Booges,  Kendall  and  Booge,  loc.  cit.    To  be  exact,  distinction  must 

be  made  between  points  determined  in  bulbs,  and  in  open  tubes. 

4  Water  decomposed  the  compounds.    The  apparatus  was  so  set  up  that 

when  it  was  necessary  to  maintain  a  low  temperature  during  filtra- 
tion, the  funnel  could  be  surrounded  by  a  freezing  mixture, 

15 


The  composition  was  calculated  from  the  volume  of  stand- 
ard alkali  required  to  neutralize  a  weighed  amount  of  the  com- 
pound. Check  determinations  were  run  to  preclude  the  possi- 
bility of  unremoved  acid  giving  erroneous  and  misleading  re- 
sults. 

FORMATE  SYSTEMS 
Formic  Acid. 

The  formic  acid  used  was  prepared  from  Baker  and  Adam- 
son  c.  p.  acid  by  treatment  with  boron  trioxide  to  remove  the 
water  it  contained.  The  mixture  was  distilled  in  vacuum, 
moisture  being  excluded1.  The  acid  which  was  used  froze  be- 
tween 8.35°— 8.5°,  and  was  generally  better  than  8.4° 2. 

SYSTEM  FORMIC  ACID— POTASSIUM  FORMATE 

This  system  was  investigated  partially,  by  GroschufP, 
who  determined  the  solubility  of  potassium  formate  in  formic 
acid  between  0°  and  100°C.  He  used  a  method  similar  to  that 
employed  in  this  work,  and  succeeded  in  isolating  an  acid  salt 
KCHO2-H2CO2,  which,  according  to  the  investigator,  under- 
went transition  before  it  melted.  In  view  of  the  incompleteness 
of  the  work,  it  was  thought  advisable  to  repeat  the  part  already 
done,  in  addition  to  completing  that  part  left  undone.  As  will 
be  seen  from  the  data,  the  course  adopted  was  justified.  Gros- 
chufPs  work  was  not  only  incomplete,  but  was  erroneous  as 
well. 

The  potassium  salt  was  prepared  by  dissolving  a  pure  sam- 
ple of  potassium  carbonate  in  90%  formic  acid.  After  ex- 
pelling the  CO2,  the  hydrate  of  potassium  formate4  was  crys- 
tallized from  the  solution.  This  salt  was  dehydrated  and  dried 
as  completely  as  possible  by  prolonged  heating  just  below  the 


1  Berichte  H  1709  (1881).     See  particularly,  Schlesinger  and  Martin, 

loc.  cit,  whose  apparatus  and  procedure  were  followed. 

2  Varying  values  are  given  in  the  literature,  8.5°  being  probably  correct. 

8. 43°,  Peterson,  Ber.  U  1191  (1880).  8.5°  Walden,  Trans.  Far.  Soc._6_ 
71  (1910).  8.52°,  Novak  (by  extrapolation),  Phil.  Mag.  (5)  44,  828. 
8.6°,  Schlesinger  and  Calvert,  loc.  cit. 

3  Berichte  36,  1783  (1903). 

4  Groschuff,  loc.  cit.,  who  also  discusses  the  hygroscopicity  of  the  salt, 

which  is  very  great. 

16 


25 


75 


100 


50 
Mol.  %  Base 

Fig.  I. 

(A) — Potassium  Formate — Formic  Acid,     (subtract  30°  from  the  tem- 
perature readings). 
(B) — Ammonium  Formate — Formic  Acid. 


17 


melting  point.  The  last  trace  of  water  was  removed  by  recrys- 
tallization  from  absolute  ethyl  alcohol,  and  dessication  over 
99%  H2SO4  in  vacuum. 

The  salt  obtained  melted  at  167.5°  ±  Q.5°C.1 

The  data  for  the  system  follow  ;  see  also  Fig  I  A. 

(a)  Solid  phase— H2CO2. 

%  Base  0  0.97         3.02         4.73         6.36         8.43 

Temp.  8.35         7.4  4.9  2.2       —0.9       —5.7 

%  Base  10.74        12.63        13.88        15.57 

Temp.  -12.6      -18.7      —23.8     —31.5 

(b)  Solid  phase— KHCO2-3H2CO2  (?) 

%  Base  16.32       16.52        17.44        18.10        18.88 

Temp.  —29.0      -27.3     —23.5      —21.7      -19.5 

(c)  Solid  phase— KHCO2-2H2CO2. 

%  Base  19.48        19.91        21.21        22.79 

Temp.  —19.0      —16.0       —8.0       —0.6 

(d)  Solid  phase— KHCO2-H2CO2. 

%  Base  23.04  24.14  25.98  28.56  30.41  31.97  33.74 
Temp.  —10.1  7.3  29.9  53.0  65.1  72.9  80.6 

%  Base  37.29        38.34        41.68       42.63        42.91        45,95        50.25 

Temp.  93.0         96.1        103.2        104.3        104.6        107.5        108.6 

%  Base  51.49        54.24        58.47        61.14       63.14 

Temp.  108.2        107.2        103.4        100.6         98.7 


(e)     Solid  phase— KHCO2 

%  Base  66.45       68.71        71.24       75.18        77.75        82.41        86.68 

Temp.  108.1        114.5        122.3        130.7        135.8        143.6        150.0 

%  Base  91.24      100.00 

Temp.  157.3        167.5 


1  The  salt  used  by  Groschuff  melted  at  157°C. 

18 


Analyses: 

(b)  It  was  not  practicable  to  carry  out  this  analysis,  be- 
cause of  the  low  temperature  at  which  compound  undergoes 
transition  ( — 17.5°).    The  slope  of  curve  indicates  the  probable 
correctness  of  the  composition  given. 

(c)  Analysis  was  carried  out  at  — 4°C.    0.3838  gm.  salt 
contained  0.2050  gm.  acid.     Calculates  to  67.9  mol.  %  acid. 
Theoretical  2 :  1  is  66.7%. 

(d)  1 :  1   compound  melts  at  108.6°  ±  0.1°.    Analysis  was 
unnecessary. 


SYSTEM  SODIUM  FORMATE— FORMIC  ACID 

This  system  had  also  been  partially,  and  as  the  work  herein 
reported  shows,  erroneously  examined  by  Groschuff1.  The 
salt  was  obtained  by  recrystallizing  a  pure  commercial  sample 
twice  from  water,  and  dehydrating  at  140°C  to  constant  weight. 
The  salt  gave  a  melting  point  of  255°  ±  I0.2 

The  curve  does  not  extend  beyond  approximately  160°C, 
because  of  the  decomposition  of  the  acid  at  this  point3.  This 
places  an  automatic  limit  upon  all  the  curves. 

The  data  for  the  system  follow.  The  curve  is  given  in 
Fig.  II  A. 


(a)     Solid  phase— H2CO2. 

%  Base  0.0  1.04         2.92         4.61          6.35          8.72 

Temp.  8.4  7.5  5.3  3.0  0.4        —3.8 


%  Base  10.50        12.58 

Temp.  —7.6    —12.8 


1  Loc.  cit. 

2  Groschuff  obtained  253°  as  the  m.  pt.  of  sodium  formate. 

3  Richter,    Org.    Chem.    1,266 — Lorin,    Jahresbericht    uber    Fortschritte 

Chemie,28,  515  (1876). 

19 


200 


150 


100 


50 


—50 


/ 


12.5 


25 
Mol.  %  Base 


Fig.  II. 

(A) — Sodium  Formate — Formic  Acid. 
(B) — Sodium  Acetate — Acetic  Acid. 


37.5 


50 


20 


(b)     Solid  phase— NaHCO2-2H2CO2. 

%  Base  14.52        16.45        18.18        20.69        21.13        21.86        23.23 

Temp.  —17.4  0.3          10.7         22.9         24.4         26.3         30.5 

%  Base  24.43        24.91 

Temp.  33.1          34.5 


(c)     Solid  phase— NaHCO2-H2CO2. 

%  Base  25.18        25.90        27.00       27.10        28.39        29.80        31.15 

Temp.  31.0          37.3         45.2         45.6          52.2          59.0         65.1 

%  Base  31.65        32.14 

Temp.  67.5          69.6 


.(d)     Solid  phase— NaHCO2. 

%  Base  29.80        31.65        33.40        35.71        39.10        42.79        43.47 

Temp.  45.1          63.6          81.2         99.3        118.6        135.2        137.7 


Analyses: 

(b)  0.2842  gm.   compound  contained  0.1615  gm.   acid; 
equivalent  to  0.1674  mole  acid  to  0.0888  mole  salt  or  66.06  mol. 
%  acid.    Theoretical  2 :  1  is  66.7r/< . 

(c)  0.2013  gm.   compound   contained  0.0814  gm.-  acid; 
equivalent  to  0.1770  mole  acid  to  0.1763  mole  salt ;  or  50.1  mol. 
%  acid.    Theoretical  1 :  1  is  50.0%. 

SYSTEM  LITHIUM  FORMATE— FORMIC  ACID 

Groschuff1  also  worked  with  this  system,  but  did  not  do  the 
part  of  the  curve  of  chief  theoretical  interest.  His  work  was 
extended  to  complete  the  curve  as  far  as  possible. 

Lithium  formate  was  prepared  from  a  pure  sample  of  the 
carbonate  and  90%  B.  and  A.  acid.  The  hydrate  was  crystal- 
lized from  the  solution,  dehydrated  at  100°C2,  and  the  dry  salt 


1  Groschuff,  loc.  cit. 

2  It  decomposes  to  water  and  anhydrous  salt  at  94°C — Groschuff,  loc.  cit. 

21 


recrystallized  from  alcohol,  and  dessicated  over  99%    H2SO4 
in  vacuum. 

The  data  for  the  system  are  given  below : 

(a)     Solid  phase— H2CO2. 

%  Base  0.0  1.58         3.47         5.33         7.09         8.93        10.75 

Temp.  8.4  7.0  5.2  3.2        +1.1        —1.3        —3.5 

%  Base  12.23        13.99        18.19        19.56       21.15        22.24       23.49 

Temp.  —5.6       —8.2      —14.6      —17.1      —19.8      —21.7      —23.5 

%  Base  24.33 

Temp.  —25.0 


(b)     Solid  phase— LiHCO2. 

%  Base  23.49       23.93       25.31        25.91        26.38       27.71 

Temp.  18.0         34.0         80.0         90.5         97.9        113.1 

%  Base  29.87       31.98       33.04       35.01        36.13 

Temp.  131.2        145.1        150.4        159.1        163.5 

Analysis  proved  (b)  to  be  the  neutral  salt. 


SYSTEM  AMMONIUM  FORMATE— FORMIC  ACID 

The  existence  of  an  acid  salt  of  ammonium  (equimolecular) 
was  shown  by  Groschuff1,  after  Reik2  had  failed  to  find  one. 
No  complete  examination  of  the  system  has  been  previously 
undertaken. 

The  ammonium  salt  was  prepared  by  passing  ammonia 
into  90%  acid.  The  acid  was  cooled  in  ice  until  it  was  almost 
saturated ;  it  was  then  allowed  to  warm  up  sufficiently  to  insure 
the  separation  of  the  neutral  salt.  This  salt  was  collected  on 


1  Groschuff,  Berichte,  36,  4351  (1903). 

2  Reik,  Monatschefte,23,  1033  (1902). 

22 


a  Biichner  funnel,  and  recrystallized  from  absolute  alcohol ;  it 
was  dessicated  over  99%  H2SO4  in  vacuum.  The  salt  gave  a 
melting  point  of  117.3°  ±  0.2°,  that  reported  by  Groschuff  being 
116°. 

The  data  for  the  system  are  given  below  and  in  Fig.  I  B  : 

(a)     Solid  phase— H2CO2. 

%   Base  0.0  1.53          3.70         6.28         8.19        10.14        12.43 

Temp.  8.47         7.0  4.5  0.6       —2.8       —6.9      —12.6 

%  Base  14.86        17.11        18.95 

Temp.  —19.8      —26.9"    —33.8 


(b)  Solid  phase— NH4HCO2-3H2CO2  (?). 

%  Base  20.43       21.73        23.33 

Temp.  —31.3      —30.0      —29.3 

(c)  Solid  phase— NH2HCO2-H2CO2,  two  crystalline  modifi- 
cations, less  soluble  (stable) — needles ;    more  soluble  (unsta- 
ble)— prisms. 

%  Base  23.33       24.83'       25.41        26.64        27.90 

Temp.*          —32.5      —26.0      —23.5      —18.7      —14.0 
Temp.f  —13.7        —8.7 

%  Base  30.11        33.25        36.29        39.30       41.90       42.94        44.38 

Temp.*  1.3  7.4          11.3          13.8          14.3         15.0 

Temp.f  +0.7          10.4          17.3         22.2         24.9         25.8 

(d)  Solid  phase— NH.HCCV 


%  Base 
Temp. 

44.38 
20.4 

46.06 
29.3 

47.87 

37.5 

51.88 
53.1 

59.98 
74.3 

68.23 
89.5 

74.32 
96.5 

%  Base 
Temp. 

76.20 
98.5 

83.46 
103.7 

88.41 
108.5 

91.64 
111.7 

100.00 
117.3 

*  Unstable.  f  Stable. 

1  At  the  higher  points,  upon  keeping  the  solutions  in  molten  condition 
for  a  while,  there  appeared  to  be  a  tendency  for  the  melting  points  to 
be  slightly  lowered,  making  it  difficult  to  check  points.  This  must 
be  due  to  decomposition  in  the  liquid  state,  possibly  to  formamide, 
although  this  generally  takes  place  at  a  much  higher  temperature — 
180°-230°  Beilstein  1,395. 

23 


Analyses: 

(b)  could  not  be  analyzed.    Nature  of  curve  makes  it  prob- 
able that  composition  given  is  correct. 

(c)  The  two  crystalline  modifications  have  the  same  com- 
position, 0.5846  gms.  compound  gave  0.2432  gm.  acid — equiva- 
lent to  0.5285  to  0.5413  moles  or  49.4  molecular  r/<  acid  (Theory 
1:  1  is  507r). 


SYSTEM  BARIUM  FORMATE— FORMIC  ACID 

There  has  been  no  work  on  solubilities  in  this  system  re- 
ported in  the  literature.  The  barium  salt  was  prepared  from 
the, carbonate  and  acid.  It  was  recrystallized  from  water  three 
times,  and  dried  at  140°C. 

The  data  are  given  below : 

(a)  Solid  phase— H2CO2. 

%  Base  0.00       0.91        1.74        2.30        3.73        4.67        5.12        6.95 

Temp.  8.4         7.2         6.1          5.1         2.6      +0.5      —0.3      —4.9 

(b)  Solid  phase— Ba(HCO2)2-H2CO2. 

%Base  8.52       8.86       9.23       9.83      10.03      10.75 

Temp.  9.5        15.5        19.0       24.9       26.5        31.8 

Analysis: 

(b)  0.4262  gm.  cpd.  contained  0.0733  gm.  acid ;  equiva- 
lent to  0.1593  to  0.1552  moles  acid  and  salt  respectively.  This 
calculates  to  50.66  mol.  %  acid.  Theory,  1 :  1  is  50%. 


SYSTEM  CALCIUM  FORMATE— FORMIC  ACID 

No  previous  work  in  this  system  could  be  located  in  the 
literature. 

The  calcium  salt  was  prepared  by  a  method  analogous  to 
that  used  for  barium  formate. 

The  data  are  given  below : 

24 


(a)  Solid  phase— H2CO2. 

%   Base  0.00         0.16         0.48         0.71          0.93          1.27          1.53 

Temp.  8.4  8.1  7.7  7.4  7.2  6.9  6.6 

(b)  Solid  phase— Ca(HCO2)2. 

%  Base  0.39  ,     0.57        0.83        1.10        1.26        1.35        1.54        1.61 

Temp.  128.6      100.0        79.0        61.0        49.7        45.5        35.0        30.0 

Analysis  proved  (b)  to  be  the  neutral  salt. 

It  is  seen  from  the  data  that  calcium  formate  in  formic 
acid  exhibits  retrograde  solubility.  This  is  not  surprising, 
since  the  same  phenomenon  occurs  in  aqueous  solutions  of 
many  calcium  "salts1. 

SYSTEM  MAGNESIUM  FORMATE— FORMIC  ACID 

This  system  has  not  been  previously  investigated.  The 
salt  was  prepared  by  dissolving  the  oxide  (Kahlbaum)  in  acid, 
and  dehydrating  the  crystallized  salt  (a  dihydrate)  at  110°C. 
The  resulting  salt  was  slightly  basic,  but  not  sufficiently  so  to 
have  any  effect  upon  the  results  obtained2. 

It  was  not  possible  to  obtain  a  curve  giving  the  solubility 
of  this  salt.  It  was  found  that  those  mixtures  which  yielded 
clear  solutions  would  not  crystallize  to  any  solid  phase  other 
than  formic  acid.  Furthermore,  on  heating  other  mixtures,  to 
get  a  more  concentrated  solution,  there  occurred  separation  of 
solid,  which  dissolved  very  slowly  on  cooling.  Measurements 
showed  that  the  magnesium  salt  was  soluble  up  to  a  concentra- 
tion of  0.2  molecular  per  cent,  at  room  temperature  (about 
25°C). 

SYSTEM  LEAD  FORMATE— FORMIC  ACID 

No  work  on  this  system  has  been  reported  in  the  literature. 
The  salt  was  prepared  by  double  decomposition  from  lead  ni- 


1  Washburn,    Physical    Chemistry,   pg.   354.     Lumsden,   J.    Chem.    Soc. 

(London),  81,  355  (1902). 

2  Analysis  showed  a  basicity  of  much  less  than  1%.    Since  the  salt  was 

only  very  slightly  soluble,  the  amount  of  base   (or  of  water  pro- 
duced) is  of  a  negligible  magnitude. 

25 


trate  and  sodium  formate.  The  resulting  precipitate  was  thor- 
oughly washed  and  recrystallized  twice  from  water.  It  was 
dried  at  140"C.  All  of  the  points,  the  data  for  which  are  given 
below,  were  determined  in  bulbs,  due  to  the  relatively  insoluble 
nature  of  the  salt. 

Solid  phase— Pb(HCO2)2. 

%   Base  0.21          0.30         0.42         0.51 

Temp.  20.0         73.1        109.4        124.5 

SYSTEM  ZINC  FORMATE— FORMIC  ACID 

No  previous  work  in  this  system  has  been  reported.  The 
salt  was  prepared  by  the  method  used  for  the  Ba  and  Ca  salts. 
The  solubility  was  less  than  0.1  mol.  %  at  140°C. 

SYSTEM  COPPER  FORMATE— FORMIC  ACID 

No  previous  work  has  been  reported.  The  salt  was  made 
from  a  pure  sample  of  basic  carbonate  and  acid.  The  hydrate 
was  crystallized  from  water  and  dehydrated,  at  about  80°C,  giv- 
ing the  bright  blue  neutral  salt1.  The  solubility  was  less  than 
0.1%  at  135°C. 

SYSTEM  NICKEL  FORMATE— FORMIC  ACID 

No  work  on  this  system  has  been  reported  in  the  literature. 
The  salt  was  prepared  from  the  carbonate  and  c.  p.  acid.  The 
hydrate,  which  was  crystallized  from  water,  was  dehydrated 
at  140°C.  It  was  soluble  to  less  than  0.1  molecular  %  at  140°C. 

SILVER  FORMATE 

Inasmuch  as  it  was  desired  to  show  the  variation  in  com- 
pound formation  below  hydrogen  in  the  electromotive  series, 
attempts  were  made  to  investigate  the  silver  system.  The  salt 
is  described  in  the  literature2  as  a  white  crystalline  salt,  de- 
composed by  boiling  water  to  give  Ag  and  CO2.  The  descrip- 


1  Voss,  Liebig's  Ann  a  1  en  266.  33  (1891). 

2  Beilstein,  I,  395.    Richter,  Organische  Chemie,  1^  266. 

26 


tion  is  not  strictly  accurate.  It  was  found  impossible  to  keep 
the  salt  in  the  presence  of  water  for-  the  length  of  time  neces- 
sary to  filter  the  solution  by  suction.  The  decomposition  is  not 
due  to  light,  as  has  also  been  suggested1,  but  undoubtedly  is 
caused  by  the  aldehydic  nature  of  the  acid  itself. 

The  salt  was  isolated  by  precipitating  it  in  absolute  methyl 
alcohol.  It 'is  not,  however,  sufficiently  stable  to  work  with; 
it  decomposed  in  a  dessicator  over  H2SO4,  even  though  pro- 
tected from  the  light. 

ACETATE  SYSTEMS 

Acetic  Acid :  100%  acetic  acid  was  prepared  from  glacial 
acetic  acid,  by  using  the  method  of  Gross2.  From  the  freezing- 
point  of  the  acid,  and  DeVisser's3  figures,  the  percentage  of 
water  was  calculated,  and  the  amount  of  acetic  anhydride  re- 
quired to  react  with  that  amount  added  to  the  acid.  The  mix- 
ture was  refluxed  for  about  30  hours  and  then  distilled. 

The  acid  used  froze  between  16.5°  and  16. 6°. 4 

SYSTEM  SODIUM  ACETATE— ACETIC  ACID 

This  system  has  not  been  previously  examined  with  re- 
spect to  solubility.  The  literature  mentions  the  existence  of  the 
acid  salts  NaAc-2HAc  melting  at  80°  and  the  equimolecular 
compound  melting  "above  140°C".5  The  "cryohydrate"  of  this 
system  has  been  studied6,  i.  e.,  the  location  of  the  eutectic  and 
the  composition  of  the  solid  phase  on  either  side  were  deter- 
mined. 

The  salt  used  was  a  pure  Baker  and  Adamson  hydrate, 
which  was  recrystallized  from  water,  and  maintained  at  140° 
for  over  a  week.  The  resulting  salt  gave  no  evidence  of  even  a 
trace  of  water. 


1  Richter,  loc.  cit. 

2  Dissertation,  Col.  University,  1919,  pg.  15. 
3DeVisser,  Rec.  Trav.  Chim.  Belg.,  12,  101  (1893). 

4  DeVisser,  loc.  cit.,  obtained  16.67  hydrogen,  scale,  equivalent  to  16.6° 

mercury  scale. 

5  See  table  I. 

6  Vasiliev,  J.  Russ.  Phys.  Chem.  Soc.,  41,  753-7  (1909). 

27 


The  data  for  the  system  are  given  below  and  in  Fig.  II  B  : 
(a)     Solid  phase— H4C2O2. 


%  Base            0.0           0.83         3.59         5.40 
Temp.               16.5         16.1          14.3         13.1 

(b)     Solid  phase—  NaH3C2O2-2H4C2O2. 

• 

%  Base             7.11          8.92        12.17        15.27        16.58 
Temp.               25.3         36.7         54.3         66.9         71.9 

21.55 
85.7 

%  Base           30.72       33.03       33.16 
Temp.                 9.61        96.25        96.3 

(c)     Solid  phase—  NaH3C202-H4C202. 

%  Base           34.03       36.87       39.06       42.54       44.25 
Temp.             112.0        132.3        145.2        157.0        160.6 

46.28 
162.3 

(d)     Solid  phase—  NaH3C2O2. 

%  Base           48.76        49.49 
Temp.              174.4        195.5 

26.86 
93.2 


SYSTEM  ZINC  ACETATE— ACETIC  ACID 

No  previous  work  in  this  system  has  been  located  in  the 
literature.  Zinc  acetate  (Kahlbaum)  was  dehydrated  at  100"C. 
The  anhydrous  salt  was  only  very  slightly  soluble  in  acetic 
acid— being  soluble  to  0.1  mol.  %  at  130°C. 

SYSTEM  FERRIC  ACETATE— ACETIC  ACID 

Neutral,  anhydrous  ferric  acetate  cannot  be  prepared.  The 
salt  used  was  basic,  probably  as  slightly  so  as  it  is  possible  to 
obtain  it.  It  was  prepared1  by  treating  a  solution  of  ferric 
acetate  in  about  90%  acetic  acid,  with  an  excess  of  acetic  anhy- 
dride, and  refluxing  the  mixture.  The  resulting  crystals  were 
dried  with  ether.  The  value  obtained  in  this  experiment  is 


1  The  salt  was  part  of  a  sample  made  by  Professor  J.  E.  Zanetti,  of  Co- 
lumbia, to  whom  the  author  wishes  to  express  his  thanks. 

28 


the  maximum  solubility  that  ferric  acetate  could  have,  since 
the  neutral  salt  would  probably  be  less  soluble  than  the  basic 
salt  used. 

The  salt  was  soluble  to  less  than  0.07%  at  140°C. 

SYSTEM  NICKEL  ACETATE— ACETIC  ACID 

There  has  been  no  work  on  this  system  reported  in  the 
literature.  The  salt  was  prepared  by  dissolving  the  carbonate 
in  c.  p.  acid,  and  crystallizing-  the  hydrated  acetate  from  the 
solution.  This  compound  was  dehydrated  and  treated  with 
acetic  anhydride  to  prevent  the  formation  of  any  basic  salt. 

The  salt  was  soluble  to  0.44%  at  140°C. 

SYSTEM  SILVER  ACETATE— ACETIC  ACID 

No  previous  work  has  been  reported.  Silver  acetate  was 
prepared  from  silver  nitrate  and  sodium  acetate.  The  precipi- 
tate was  thoroughly  washed,  and  recrystallized  from  water.  It 
was  dried  over  99%  H2SO4  in  vacuum.  The  measurements 
could  not  be  carried  very  high  (not  above  115°),  as  the  acetate 
underwent  reduction  at  higher  temperature.  The  silver  salt  is 
not  very  soluble  in  acetic  acid.  The  data  are  given  below : 

%   Base  0.094        0.204 

Temp..  76.0        115,0 

CONDUCTIVITY  MEASUREMENTS 

To  complete  the  series  of  conductivities  compiled  by 
Schlesinger  and  his  collaborators1,  measurements  were  made 
of  the  conductivity  of  systems  Ba,  Pb,  and  Mg  formates  in 
formic  acid.  The  measurements  were  made  in  a  cell  of  the 
Freas  type,  with  platinized  electrodes.  Inasmuch  as  the  preci- 
sion required  for  a  qualitative  comparison  of  conductivities  was 
not  very  high,  it  was  not  found  necessary  to  balance  out  the 
capacity  of  the  cell  by  means  of  a  condenser.  The  results  are 
accurate  to  better  than  —  1%. 


1  Loc.  cit. 

29 


Conductivities  were  determined  in  a  thermostat  main- 
tained at  25.00°  ±  0.01°.  A  Leeds  and  Northrup  bridge  of  the 
Kohlrausch  type,  with  telephone  receiver  tuned  to  1000  cycles 
was  used.  The  current  was  supplied  by  a  constant  speed  high 
frequency  generator  (1000  cycles/sec.). 

No  attempt  was  made  to  get  formic  acid  of  conductivity 
as  low  as  that  obtained  by  Schlesinger  and  his  co-workers1. 
That  used  in  this  work  had  a  specific  conductivity  of  7.3 — 
7.5  X  10~r>  reciprocal  ohms. 

CONDUCTIVITIES 

34  44 

Base  Cone.2  S  X  10  Sa  X  10 

Barium  0.2758  11.92  11.85 

0.1869  8.551  8.478 

0.0923  4.902  4.827 

0.0474  2.649  2.575 

Lead  0.1047  2.462  2.387 

0.0477  1.370  1.295 

Magnesium  0.0919  3.006  2.931 

0.0479  1.758  1.683 

DISCUSSION  OF  RESULTS 

An  examination  of  the  experimental  results  will  disclose 
the  extent  of  agreement  between  them  and  the  corresponding 
consequences  of  the  hypothesis  proposed  in  the  beginning  of 
the  paper. 

(a)     Compounds  Isolated. 

The  curves  in  general,  and  the  particular  compounds  iso- 
lated may  be  first  compared.  The  addition  compounds  crystal- 
lized, generally,  as  needles,  whereas  the  neutral  salts  give  in 
most  cases  crystals  belonging  to  the  rhombic  system.  Table 
II  gives  the  compounds  actually  isolated,  with  the  freezing- 
point  relationships  of  each. 


1  S.  and  Martin,  J.  A.  C.  S.,  36,  1590  (1914). 

2  Cone,  in  equivalents  per  liter. 

3  S=Specific  conductivity  of  solution  in  reciprocal  ohms. 

4  Sa=S — Specific  conductivity  of  solvent. 

30 


Compound 
KHC02-3H2C02 1 
KHCO^H.CCV 
KHC02-H2C02 
NH4HCO2-3H2CCV 
NH,HC02-H2C022 

NaHCO2-2H2C(V 

NaHCO2-H2CO2 

Ba(HC02)2-H2C021 


TABLE   II. 


Transition  Temp. 
—18.0 
4.0 
Melts  at  108.6° 

—29.0 

prisms  27.0 

needles  14.8 

35.6 

70.5 

undeterminable" 


Composition  solution 

in  equilibrium  at 
trans,  temp.  (moL  %) 

19.6 
23.9 

23.3 
45.6 
43.1 

25.7 
32.3 


The  order  of  increasing  complexity,  as  well  as  of  increa's- 
ing  stability  of  the  complexes,  is  seen  to  be  that  required  by  the 
hypothesis.  Potassium  forms  the  most  complex  compounds, 
and  of  the  eight  different  compounds  isolated,  the  equimolecu- 
lar  potassium  acid  formate  is  the  only  one  sufficiently  stable  to 
give  an  actual  melting  point4.  The  form  of  the  curve  (Fig. 
I  A)  in  the  neighborhood  of  the  maximum  indicates  some  dis- 
sociation of  the  compound  into  its'  components5.  The  more 
complex  components  undergo  transition  before  they  melt. 

Next  in  order  of  complexity  and  stability  to  the  potassium 
compounds,  are  those  of  ammonium  formate  (Fig.  IB).  A 
comparison  of  the  extent  of  compound  formation6  shows  sod- 
ium and  ammonium  formates  to  be  solvated  (in  solution)  to 
practically  the  same  extent.  The  increased  complexity  and 
stability  in  the  case  of  the  ammonium  compounds,  is  undoubt-' 
edly  due  to  the  temperature  factor7 — the  lower  temperature  at 
which  the  ammonium  complexes  exist,  decreases  their  tendency 


1  Compounds  not  previously  mentioned  in  the  literature. 

2  Two  modifications  not  previously  noted. 

3  See  below. 

4  Groschuff,  loc.  cit,  did  not  obtain  a  melting-point,  as  noted. above. 

5  The  relation  between  stability  of  the  compound  and  the  sharpness  of 

the  maximum  is  discussed  by  Kendall  and  Booge,  loc.  cit,  pg.  1728; 
also  by  Kremann,  Monatshefte_25,  1215  .(1904). 

6  See  section  (b),  following. 

7  See  particularly,  Gross,  loc.  cit.,  p.  29. 


31 


to  decompose.  It  is  interesting  that  the  equimolecular  com- 
pound exists  in  two  crystalline  forms,  the  stable  form  being 
the  only  one  previously  mentioned1. 

The  sodium  compounds  are  quite  unstable,  undergoing 
transition  into  compounds  of  a  lower  order  of  complexity  long 
before  their  respective  melting  points  are  reached. 

The  same  is  true  of  the  barium  compound,  the  transition 
point  of  which  is  given  as  indeterminable.  The  last  few  points 
on  the  curve  are  probably  metastable.  On  standing,  crystals 
separate  which  are  probably  the  anacid  salt.  Not  enough  of 
these  could  be  obtained  for  an  analysis,  nor  could  they  be  made 
to  separate  in  fine  enough  form  to  enable  one  to  determine  a 
melting-point,  i.  e.,  to  locate  the  stable  curve.  Solutions  more 
concentrated  than  the  last  one  could  not  be  obtained  at  140"C. 

Lithium  formate,  though  soluble  in  formic  acid,  does  not 
form  any  isolable  complex  with  it.  It  is  noteworthy  that  this 
base,  that  of  the  least  electropositive  of  the  alkali  metals,  yields 
no  compound,  whereas  barium,  the  most  electropositive  of  the 
common  alkaline  earths,  forms  an  equimolecular  compound 
with  formic  acid.  Other  properties  also  place  these  two  metals 
in  this  order2. 

(b)  Extent  of  compound  formation  from  the  relative  slopes  of  the  curves. 

It  has  already  been  emphasized  that  addition  compounds 
may  be  formed  in  solution,  and  yet  not  be  sufficiently  stable  to 
allow  of  their  being  isolated3.  By  applying  a  more  general 
method  of  detecting  compound  formation,  it  was  shown  that*  a 
better  estimate  of  relative  degrees  of  solvation  in  solution  was 
obtainable.  This  method  consists  in  the  determination  of  the 
extent  of  deviation  of  the  curve  representing  the  data  collected, 
from  the  ideal  curve4,  the  equation  of  which  is 


1  Groschuff,  loc.  cit,  who  gives  the  transition  temperature  as  23.5°,  in- 
stead of  27.0°  obtained  in  this  work. 

2Abegg,  Hand.  Anorg.  Chem.  2-1,  117.  Soddy,  Chemistry  of  the  Radio 
Elements,  p.  44,  gives  the  order  of  the  elements  on  the  basis  of 
their  physical  properties  as  K,  Na,  Ba,  Sr,  Li,  Ca. 

3  Kendall,  J.  A.  C.  S,  36,  1731  (1914).    Kendall  and  Booge,  loc.  cit.,  p. 

1730,  Kendall,  Booge  and  Andrews,  loc.  cit.,  2309. 

4  Washburn,  Phys.  Chem.  p.  174,  Kendall  and  Booge^ibid. 

32 


+10 


-10 


—20 


-30 


—40 


10 
Mol.  %  Base 


15 


20 


Fig.  III. 

Freezing-point  depressions.    In  order  of  increasing  depression : 
a — Ideal      b — Lithium      c — Sodium      d — Ammonium      c — Potassium. 


33 


.          -.OgeX=^ 

The  factors  affecting  the  precision  with  which  an  estimate 
of  the  degree  of  hydration  of  electrolytes  in  aqueous  solution 
can  be  made,  are  fully  discussed  in  the  articles  to  which  refer- 
ences has  been  made2.  It  has  been  demonstrated  that  a  quanti- 
tative estimate  of  solvation  is  at  present  impossible.  In  a  series 
such  as  that  being  discussed,  the  relative  extents  of  solvation 
will  be  given  by  the  respective  deviations  from  the  ideal  curve. 

The  graphs  of  the  ideal  and  observed  curves  are  given  in 
Fig.  III.  As  no  substance  was  obtainable  which  would  give  an 
ideal  curve  with  formic  acid,  the  ideal  curve  was  calculated. 
The  known  values  of  Q3  and  of  T0  4  were  substituted  in  the 
equation  already  given,  and  the  value  of  T  corresponding  to 
a  given  value  of  X  determined.  The  data  from  which  the  ideal 
curve  was  plotted,  follow  (solid  phase  is  H2CO2  throughout) : 

Mol.  fract.  Base        0.00       0.02      0.05      0.11      0.14      0.18      0.22      0.08 
Temp.  281.43    280.1    278.1    274.0    271.9    269.1    266.1    276.1 

The  curves  indicate  that  the  extent  of  divergence  increases 
in  the  order  Li,  Na,  NH4,  K,  in  agreement  with  the  theory.  The 
bivalent  bases  of  Ca  and  Ba  lie  below  that  of  potassium  in  the 
order  named5.  They  are  not  comparable  with  the  alkali  metal 
bases,  because  the  nature  of  their  ionization  is  unknown0.  It  is 
probable  that  three  ions  are^ formed  to  some  extent.  It  is  evi- 
dent, however,  that  barium  formate  is  more  extensively  sol- 
vated  than  calcium  formate. 


1  X  =  Mol.  fraction  solvent ;  Q,  molar  heat  of  fusion  of  the  "solvent." 

R=  1.988  cal.    (gas   constant).     TQ   is   temperature   (absolute)    of 
fusion  of  solvent  and  T  that  of  the  solution. 

2  The  factors  considered  by  Kendall,  Booge  and  Andrews,  loc.  cit.,  ob- 

tain in  the  system  under  consideration,  and  the  conclusions  arrived 
at  in  that  paper  are  equally  significant  here. 

3  Q  — 2421.2,  Berthelot,  C.  R.  78,  716  (1874). 
*Tft  =  281.43  Peterssen,  Ber.  13,  1191  (1880). 

5Ca  and  Ba  are  not  given  in  Fig.  Ill,  at  1.5%  the  f.  pt.  is  6.6°;  Ba  at 

1.5%  is  at  6.3°. 
6  Schlesinger  and  Mullinix,  loc.  cit.,  pg.  75.     Schlesinger  and  Bunting, 

J.  A.  C.  S.  4M945  (1919). 

34 


The  divergence  which  has  been  noted  results  essentially 
from  ionization  and  from  solvation.  lonization  increases  the 
number  of  moles  of  solute  and  hence  produces  abnormal  de- 
pressions of  the  freezing-point.  Solvation  removes  solvent, 
thereby  increasing  the  molecular  fraction  (1 — X)  of  solute, 
and  consequently  has  a  similar  effect.  It  is  not  possible  to 
determine  exactly  what  part  of  the  total  effect  is  produced  by 
each  of  these  factors.  In  view  of  the  fact  that  it  has  been 
definitely  established  that  ''ionization  is  preceded  by  combina- 
tion between  solvent  and  solute,  and  is  indeed  a  consequence  of 
such  combination"  1,  such  an  attempt  appears  to  be  superfluous. 
It  is  noteworthy,  however,  that  the  alkali  bases,  which  are 
ionized  to  practically  the  same  extent2,  give  curves  which  are 
quite  distinct  showing  that  the  second  factor,  solvation,  is  op- 
erative. 

(c)     Solubility. 

In  addition  to  those  bases  which  have  been  considered  and 
classified  according  to  the  extent  of  compound  formation,  there 
remain  those  bases  which  were  not  sufficiently  soluble  to  allow 
a  determination  of  relative  degrees  of  solvation  by  either  of  the 
methods  described.  For  these  it  is  necessary  to  use  solubility 
as  the  criterion,  as  will  be  evident  from  what  follows. 

In  an  ideal  binary  system,  the  equation  given  above3  repre- 
sents the  equation  of  the  solubility  curve  of  either  substance  in 
the  other,  according  as  the  values  substituted  for  Q  and  T0  be- 
long to  one  component  or  the  other.  If,  then,  solubility  meas- 
urements in  any  solvent  are  made  on  a  series  of  salts  having 
approximately  the  same  values  of  Q  and  T0,  the  solubility  at 
any  given  temperature  will  be  the  same  in  all  cases  where  the 
solutions  resulting  are  ideal. 


1  Kendall  and  Booge,  loc.  cit,  p.  2324. 

2  Schlesinger  et  al,  loc.  cit. 

3  See,  particularly,  Washburn  and  Read,  Proc.  Nat.  Acad.  Sci.  J[,  191 

-    (1915).     C.  A.,  9,  1520  (1915).     Washburn,  Principles  of  Physical 
Chemistry,  pg.  172. 

35 


If,  however,  there  occurs  combination  between  the  solvent 
and  some  of  the  substances,  the  curves  for  these  will  fall  away 
from  the  maximum  more  rapidly  than  the  ideal1.  '  The  sub- 
stances will  be,  at  any  given  temperature,  more  soluble  than 
those  which  give  ideal  solutions. 

In  actual  practice,  a  series  of  salts  in  which  the  values  of 
Q  and  T0  are  the  same  throughout,  is  never  encountered.  This 
fact  removes  the  possibility  of  obtaining  any  quantitatively 
comparable  data.  In  the  case  of  the  series  under  consideration, 
the  values  of  T02  are  several  hundred  degrees  above  room  tem- 
perature. Because  of  the  large  value  of  T0,  the  solubility  will 
of  necessity  be  very  small  in  all  systems  in  which  no  compound 
formation  occurs.  The  order  of  solubility  will  accordingly 
correspond  qualitatively  with  the  extent  to  which  combination 
between  the  two  components  of  the  system  takes  place. 

The  greater  the  extent  of  compound  formation,  the  greater 
the  solubility  ought  to  be  (provided  the  comparison  be  made 
of  analogous  compounds  at  the  same  temperature).  Since  it 
has  been  postulated  that  compound  formation  should  vary  with 
the  position  of  the  metal  in  the  electromotive  series,  it  follows 
that  a  similar  variation  in  solubility  should  occur. 

Examination  of  the  data  (also  Figures  I  and  II)  shows 
that  the  solubility,  at  25°  say,  does  decrease  to  a  minimum  from 
potassium  through  the  alkaline  earths  to  zinc  and  copper.  Am- 
monium formate  is  far  more  soluble  than  potassium  formate, 
but  this  is  undoubtedly  due  to  the  low  value  of  T0  (300°  abso- 
lute for  the  equimolecular  compound).  The  only  real  excep- 
tion is  in  the  case  of  the  lead  salt,  which  is  too  soluble.  While 
no  explanation  can  be  offered  at  present,  it  is  significant  that  in 
sulfuric  acid3,  and  in  water,  the  corresponding  lead  salts  are 
also  out  of  their  proper  place. 


1  This  is  evident  from  Fig.  III. 

-  The  values  of  TQ  are  not  determinable,  except  in  the  cases  of  Na,  NH4 

and  K.     The  melting-points  of  K  and  NH4  are  low  (see  above). 

The  temperatures  of  decomposition  are  generally  above  300°C.    See 

Berichte,  51.  399  (1918). 
3  Unpublished  work  for  Davidson,  who  discusses  the  subject  of  solubility 

in  greater  detail. 

36 


It  has  already  been  mentioned  that  it  was  not  possible  to 
demonstrate  with  formates  alone  that  the  solubility  of  the  bases 
passed  through  a  minimum  and  then,  with  increasing  diverg- 
ence between  H  and  the  metal  (whose  position  is  below  hydro- 
gen), increases  again.  The  supplementary  experiments  with 
the  acetates  are  more  satisfactory  in  this  respect. 

Sodium  acetate  (Fig.  II  B)  is  very  soluble  in  acetic  acid; 
its  solubility  is  of  the  same  order  of  magnitude  as  that  of  sod- 
ium formate  in  formic  acid.  Two  compounds  were  isolated. 
The  compound  NaH3C2O2-2H4C2O2  undergoes  transition  at  its 
melting  point,  96.3°  ±0.1°  (the  literature  gives  SQ°C.)1.  The 
equimolecular  compound  undergoes  transition  at  163°,  before 
reaching  its  melting  point.  The  curve  resembles  very  much 
that  for  sodium  sulfate  in  sulfuric  acid2. 

Zinc  acetate  is  soluble  to  only  a  slight  extent,  0.1%  at 
130°C,  and  the  results  obtained  with  a  slightly  basic  ferric  ace- 
tate indicate  a  still  smaller  solubility  in  the  case  of  the  neutral 
ferric  salt.  Silver  acetate  is  several  times  as  soluble  as  the  ace- 
tates of  these  metals.  The  acetate  of  nickel  is  abnormally  solu- 
ble, 0.44%  at  140°,  for  which  fact  no  immediate  explanation  is 
available. 

The  experimental  results  taken  collectively  indicate  that 
compound  formation  and  the  related  properties,  decrease  to  a 
minimum  and  then  increase  again,  when  the  metal  (R)  in  the 
system  HX-RX  is  varied  from  the  upper  to  the  lower  end  of 
the  electromotive  series. 

While  the  increase  in  compound  formation  below  hydro- 
gen in  the  acetate  series  is  not  as  large  as  might  be  expected 
(for  example,  it  is  very  striking  in  the  sulfate  series),  it  is 
doubly  significant,  because  it  brings  up  for  consideration  a 
point  which  might  otherwise  be  overlooked.  On  comparing  the 
data  available  in  the  different  series— RSO4  in  H2SO4,  RHCO2 
in  H2CO2,  etc.,  it  becomes  evident  that  while  the  alkali  bases 
ire  soluble  to  practically  the  same  extent  in  each  series,  the  rate 
at  which  the  solubility  falls  off  as  R  is  varied  through  the  alka- 
line eartns  toward  hydrogen,  is  different  in  the  several  series. 


1  Lescoeurs,  loc.  cit. 

2  Landon,  loc.  cit.,  pg.  23. 


37 


Thus,  the  solubility  of  the  sulfates  falls  off  less  rapidly  than 
that  of  the  formates,  the  order  being  RSO4  <  RHCCX  < 
RH3C2O2<ROH1.  There  is  an  inverse  variation  in  the  extent 
to  which  the  solubility  increases  in  the  case  of  the  bases  of  the 
more  noble  metals.  Here  the  sulfates  show  the  most  pro- 
nounced increase  in  solubility,  and  the  increase  in  the  acetate 
and  hydroxide  series  is  small.  The  complete  significance  of 
these  facts  and  their  proper  explanation  may  be  evident  after 
several  series  of  fairly  strong  acids  have  been  examined,  and 
the  influence  of  X  completely  determined. 

(d)     Discussion  of  Conductivity  Measurements. 

TABLE   III. 


Base 

Cpds.  Isolated 

Cone.2 

A3 

Cone.2 

A' 

NH. 

1:3;  1:1 

*0.04646 

67.93 

*0.09293 

65.74 

K 

1:3;  1:2;  1:1 

'0.05480 

65.74 

*0.09464 

63.37 

Na 

1:2;  1:1 

*0.04381 

62.19 

fO.0923 

60.01 

Li 

none 

*0.9574 

58.79 

*0.0902 

56.68 

Ba 

1:1 

0.0474 

54.32 

0.0923 

52.30 

Sr 

not  examined 

*0.0476 

53.57 

tO.0923 

49.34 

Ca 

none 

fO.0474 

51.37 

fO.0923 

46.66 

Mg 

none 

0.0479 

35.14 

0.0919 

31.90 

Pb 

none 

0.0477 

27.15 

0.1047 

22.74 

From  this  table  it  is  evident  that  the  variation  of  compound 
formation  with  the'  position  of  the  metal  in  the  electromotive 


1  To  test  this  statement,  a  few  measurements  on  calcium  acetate  in 
acetic1  acid  were  made.  At  30°,  calcium  acetate  is  soluble  to  less 
than  0.35%  ;  calcium  formate  is  soluble  to  1.6%  ;  calcium  sulfate 
(or  rather,  a  1:3  compound)  to  5.4%,  and  calcium  hydroxide  (at 
20°)  to  0.04%,  each  in  the  corresponding  solvent.  Calcium  for- 
mate, it  has  been  noted,  shows  retrograde  solubility,  and  therefore 
the  difference  in  solubility  between  it  and  calcium  acetate  would  be 
less  at  higher  temperatures. 

-  Cone,  in  equivalents  per  liter. 

3  A  is  equivalent  conductivity  in  reciprocal  ohms. 

*  Schlesinger  and  collaborators'  results,  loc.  cit. 

f  Results  marked  thus  f,  are  calculated  by  interpolation  from  the  rate 
of  change  of  A  with  the  concentration  in  the  neighborhood  of  this 
value  of  the  concentration.  It  does  not  change  the  order  of  the 
bases,  although  necessarily  inexact.  Also,  the  difference  in  concen- 
tration in  the  different  cases  to  be  compared,  is  too  small  to  affect 
the  order  of  the  bases. 

38 


series  is  paralleled  by  the  change  in  conductivity.  The  agree- 
ment between  the  freezing-point  measurements  and  conduc- 
tivity measurements  is  good. 

It  is  again  difficult  to  fix  exactly  the  correct  positions  of 
lithium  and  barium.  Lithium  has  a  higher  equivalent  conduc- 
tance than  barium,  although  the  latter  gives  a  stable  compound 
with  formic  acid  while  lithium  does  not.  The  light  lithium  ion 
may  be  more  mobile  than  the  barium  ion. 

The  exact  position  of  ammonium  formate  among  the  bases 
is  not  easily  located.  Although  less  extensively  solvated  than 
potassium,  it  shows  greater  conductivity.  This  is  probably 
explained  by  the  fact  that  the  ammonium  complexes  are  formed 
are  less  stable1,  and  therefore  undergo  dissociation  into  ions  to 
a  greater  extent  than  do  the  analogous  potassium  complexes. 

Lead,  in  the  conductivity  results  also,  takes  the  same  ab- 
normal position  as  before.  Its  solubility  was  slightly  lower 
than  that  of  magnesium,  and  the  conductivity  results  agree 
with  this  fact. 

The  results  indicate  that  the  high  conductivity  of  formates 
in  formic  acid  noted  by  Schlesinger  and  his  collaborators,  is 
due,  as  predicted  by  the  generalization,  to  extensive  combina- 
tion between  solvent  and  solute.  The  extent  of  combination  de- 
creases in  order  from  potassium  through  the  alkali  metals  and 
alkaline  earths  to  lead,  and  is  paralleled  practically  exactly  by 
the  diminution  of  conductivity. 

The  agreement  of  the  experimental  results  with  the  argu- 
ment advanced  at  the  beginning  of  the  paper,  affords  consider- 
able support  to  the  general  validity  of  the  propositions  there 
stated. 

SUMMARY: 

The  attempt  has  been  made  to  extend  the  generalization 
correlating  compound  formation  and  chemical  diversity  to  a 
series  of  non-aqueous  systems  of  the  type  HX-RX.  On  the 
basis  of  the  generalization,  the  extent  of  compound  formation, 


1  Gross  (dissertation,  pg.  7),  has  shown  that  the  extent  of  ionization  de- 
pends upon  :  (1)  The  extent  of  the  compound  formation — AB  +  CD 
^  AB.CD  and  (b)  The  extent  of  dissociation  of  the  complexes  into 
ions  of  opposite  charge  AB .  CD  ^  (AB .  C)+  +  D-. 

39 


and  hence  of  ionization,  should  vary  with  the  difference  in  char- 
acter of  the  two  components. 

The  use  of  the  relative  positions  of  the  metals  and  hydro- 
gen in  the  electromotive  series  as  the  criterion  of  diversity  has 
been  proposed.  From  this  it  follows  that  as  R  is  varied  from 
one  extreme  of  the  series  to  the  other,  the  extent  of  compound 
formation  should  diminish  to  a  minimum  (near  hydrogen)  and 
then  increase  again. 

This  conclusion  has  been  tested  experimentally  by  the 
determination  of  freezing-point  curves  of  representative  for- 
mate-formic acid  and  acetate-acetic  acid  systems.  In  addition, 
conductivity  measurements  have  been  employed  to  test  the 
validity  of  the  argument  advanced. 

The  results  have  agreed  strikingly  with  the  deductions 
from  the  fundamental  assumptions.  Five  new  acid  formates 
were  isolated.  In  the  formate  series,  the  extent  of  combina- 
tion between  the  two  components  decreases  to  a  minimum  as 
the  position  of  R  is  varied  from  potassium  through  the  series 
toward  hydrogen.  The  acetate  series  exhibits  a  similar  varia- 
tion, with  a  slight  increase  in  the  extent  of  combination  in  the 
case  of  the  silver  system.  The  change  in  conductivity  and  in 
solubility  parallel  the  variation  in  extent  of  combination.  This 
experimental  evidence  indicates  the  probable  validity  of  the  as- 
sumptions to  which  reference  has  been  made. 


40 


VITA. 

Howard  Adler  was  born  in  New  York  City  on  January  9, 
1896,  and  attended  the  grade  and  high  schools  of  that  city.  He 
received  the  degree  of  B.  S.  from  the  College  of  the  City  of 
New  York  in  February,  1916.  From  1916  to  1917  he  taught  in 
the  Department  of  Chemistry  of  the  College  of  the  City  of  New 
York,  at  the  same  time  attending  the  graduate  school  of  Colum- 
bia University.  In  June,  1917,  he  received  from  Columbia  the 
degree  of  M.  A.  From  September,  1917,  until  February,  1919, 
he  was  in  the  United  States  Army.  In  February,  1919,  he  re- 
sumed graduate  work  at  Columbia  University. 


THIS 


AN  INITIAL  FINE  OF  25  CENTS 


OVERDUE. 


7"    300ct'56vv; 


LD  21-95m-7,'37 


(iaylord  Bros. 

Makers 
Syracuse,  N. 
PH.  W.  21, 1908 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


